Types of Chemical Reactions Class 10 Notes | EduRev

CHEMICAL REACTIONS
The reaction in which a chemical substance transforms into another new chemical substance is known as a chemical reaction.

A chemical change involves the formation of new substances whereas a physical change involves a change in colour or state and no new substances are formed.

Accordingly, the reactions are classified in different types:
1. Combination reaction/Synthesis reaction
2. Decomposition reaction/Analysis reaction
3. Displacement reaction/Substitution reaction
4. Double displacement reaction
5. Oxidation and Reduction reaction
6. Endothermic and Exothermic reactions

During a chemical reaction, the substances that react are known as Reactants whereas the substances that are formed during a chemical reaction are known as Products.

Six common types of chemical reactions are discussed below:

1. COMBINATION REACTION
The reactions in which two or more substances combine to form a single new substance are called combination or synthesis reactions.

Fig: Combination Reaction

For example: Burning of coal
C(s) + O₂(g) → CO₂(g)

Combination reactions are of three common types.
(a) Combination of two elements to form a compound.
Example:

Formation of water from H₂(g) and O₂(g) 

(b) Combination of an element and a compound to form a new compound.
Example:

(c) Combination of two compounds to form a new compound.
Example: 

• Formation of Ammonium chloride: Ammonium chloride is formed by combining vapours of ammonia with hydrogen chloride gas. It is a white-coloured solid.

Fig: Ammonium chloride

• Formation of Calcium Hydroxide: Reaction between quick lime (Calcium oxide, CaO) and water is a combination reaction. In this reaction, quick lime reacts with water to form slaked lime (calcium hydroxide, Ca(OH)₂).

The reaction between quick lime and water is highly vigorous as well as exothermic.

Fig: Slaked lime

2. DECOMPOSITION REACTION
The reaction in which a single compound breaks up into two or more simpler substances is known as Decomposition reaction. The decomposition reaction generally takes place when energy in some forms such as heat, electricity or light is supplied to the reactants.
The general equation that describes a decomposition reaction is:

Fig: Decomposition Reaction

Types of Decomposition Reactions
Decomposition reactions can be classified into three types:
(a) Thermal decomposition reaction
(b) Electrolytic decomposition reaction
(c) Photo decomposition reaction

(a) Decomposition reactions by heat (Thermal decomposition)
Thermal decomposition is a chemical reaction where a single substance breaks into two or more simple substances when heated. The reaction is usually endothermic because heat is required to break the bonds present in the substance.

Examples:

• Decomposition of calcium carbonate: Calcium carbonate (limestone) decomposes into calcium oxide (quick lime) and carbon dioxide when heated. Quick lime is the major constituent of cement.
  
• Decomposition of ferrous sulphate: Ferrous sulphate crystals contain water molecules (FeSO₄. 7H₂O). On heating, ferrous sulphate crystals lose water and anhydrous ferrous sulphate (FeSO₄) is formed. So their colour changes from light green to white.

Fig: Thermal Decomposition

On further heating, anhydrous ferrous sulphate decomposes to form ferric oxide (Fe₂O₃), sulphur dioxide (SO₂) and sulphur trioxide (SO₃). So, the gas emitted smells like burning sulphur.
In this reaction, the single reactant FeSO₄ decomposes to form three different products. So, the reaction is a decomposition reaction.

• Decomposition of lead nitrate: Lead nitrate decomposes on heating and gives a lead monoxide and brown fumes of nitrogen dioxide and colourless gas oxygen are liberated.
  

(b) Decomposition by electricity (Electrical decomposition or Electrolysis)
Electrolytic decomposition may result when electric current is passed through an aqueous solution of a compound. A good example is the electrolysis of water.

Fig: Electrolysis of water

Electrolysis of water: Electrolysis of water is the decomposition of water into hydrogen and oxygen due to the passage of electric current through it.

(c) Decomposition by sunlight (Photochemical decomposition)

Fig: Decomposition of silver chloride

Decomposition of silver chloride: Place a small quantity of silver chloride (AgCl) taken in a watch glass under sunlight for some time. The crystals slowly acquire a grey colour. On analysis, it is found that the sunlight has caused decomposition of silver chloride into silver and chlorine.  
  

Silver bromide also decomposes in the same way

The decomposition of a compound with light is called Photolysis. The above reactions are used in black and white photography.

Note:
1. All the decomposition reaction requires energy i.e. these reactions are Endothermic reactions. These reactions are used in Extraction of metals.
2. A decomposition reaction is called the opposite of a combination reaction. This can be supported by the following reactions:
Combination reaction:

Decomposition reaction:

3. DISPLACEMENT REACTION
The chemical reactions in which one element takes the place of another element in a compound are called Displacement Reactions.

Fig: Displacement Reaction

Example: Reaction of iron nails with copper sulphate solution.

In this reaction, iron has displaced copper from copper sulphate solution.

Q.1. How is the chemical reactivity of metals linked with their position in the electrochemical series?
Ans. Chemical reactivity of metals is linked with their relative positions in the activity series. Certain metals have the capacity to displace some metals from the aqueous solutions of their salts. A metal placed higher in the activity series can displace the metal that occupies a lower position from the aqueous solution of its salt. The displacement reaction is not limited to metals only. Even non-metals can take part in these reactions.
Example:Halogens, the activity series of halogen is F > Cl > Br > I.

Fig: Reactivity Series of metals

Other examples of displacement reactions are:
(a) Zinc displaces copper from copper sulphate

Iron, zinc and lead are more reactive elements than so they displace copper from its compounds.
(b) Chlorine displaces Bromine from Potassium bromide

(c) Copper displaces silver from silver nitrate 

Q.2. In the refining of silver, the recovery of silver from silver nitrate solution involved displacement by copper metal. Write down the reaction involved.
Ans.


4. DOUBLE DISPLACEMENT REACTION
The reactions in which two compounds react to form two different compounds by mutual exchange of ions are called double displacement reactions. The general equation which represents a double displacement reaction can be written as:

Double displacement reactions generally take place in aqueous solutions in which the ions precipitate and there is an exchange of ions.

Fig: Example of Double Displacement Reaction

Example: On mixing a solution of barium chloride with sodium sulphate, a white precipitate of barium sulphate is immediately formed. These reactions are ionic in nature. The reactants change into ions when dissolved in water and there is an exchange of ions in solution. This results in the formation of product molecules.

Two common types of double displacement reactions are:
(a) Precipitation reaction

• A chemical reaction that involves the formation of an insoluble product (precipitate; solid) is called Precipitation reaction.
• The reactants are soluble, but the product formed would be insoluble and separates out as a solid.
• The chemical equation by which a chemical change is described is adequate for reaction in solutions, but for reactions of ionic compounds in aqueous solution (water), the typical molecular equation has different representations.
• A molecular equation may indicate formulas of reactants and products that are not present and eliminate completely the formulas of the ions that are the real reactants and products.
• If the substance in the molecular equation that is actually present as dissociated ions are written in the form of their ions, the result is an ionic equation.

A precipitation reaction occurs when a solution, originally containing dissolved species, produces a solid, which generally is denser and falls to the bottom of the reaction vessel. The most common precipitation reactions occurring in aqueous solution involve the formation of an insoluble ionic compound when two solutions containing soluble compounds are mixed. 

Consider what happens when an aqueous solution of NaCl is added to an aqueous solution of AgNO₃. The first solution contains hydrated Na⁺ and Cl⁻ ions and the second solution, Ag⁺, and NO₃⁻ ions.

NaCl(s) → Na⁺(aq) + Cl⁻(aq)
AgNO₃(s) → Ag⁺(aq) + NO3⁻(aq)

When mixed, a double displacement reaction takes place, forming the soluble compound NaNO₃ and the insoluble compound AgCl. In the reaction vessel the Ag⁺ and Cl⁻ ions combine, and a white solid precipitated from the solution. As the solid precipitates, the Na⁺ and NO₃⁻ ions remain in solution.
The overall double displacement reaction is represented by the following balanced equation:

NaCl(aq) + AgNO₃(aq) → AgCl(s) + NaNO₃(aq)

Examples 1. 

Fig: Precipitation reaction

Example 2.
(i) When aqueous solutions of Pb(NO₃)₂ and KI are mixed, does a precipitate form?
(ii) Write a balanced equation for the precipitation reaction that occurs when aqueous solutions of copper(II) iodide and potassium hydroxide are combined.
Solution. You are asked to predict whether a precipitate will form during a chemical reaction and to write a balanced equation for a precipitation reaction. You are given the identity of two reactants.
(i) Yes, a solid precipitate, PbI₂, forms when these solutions are mixed:
Pb(NO₃)₂(aq) + KI(aq) → PbI₂(s) + 2KNO₃(aq)
(ii) The two products of the reaction are insoluble copper (II) hydroxide and soluble potassium iodide.
CuI₂(aq) + 2 KOH(aq) → Cu(OH)₂(s) + 2 KI(aq)

(b) Neutralisation reaction
When an acid reacts with a base to form a salt and water by exchange of ions.
Example:

5. OXIDATION AND REDUCTION REACTIONS (REDOX REACTION)
Fig: Redox ReactionOxidation:
(i) The addition of oxygen to an element or compound

(ii) Removal of hydrogen from a compound is known as oxidation.

Reduction:
(i) The addition of hydrogen to an element or compound

(ii) Removal of oxygen from a compound.

Oxidising agent: The substance which gives oxygen or removes hydrogen for oxidation is called oxidising agent and the substance which gains oxygen during reaction is said to be oxidised.

Reducing agent: The substance which gives hydrogen or removes oxygen for reduction is called the reducing agent. The substance which gains hydrogen during reaction is said to be reduced.

Fig: Electron movement in Oxidising and Reducing agent

Those reactions in which oxidation and reduction (both) occur simultaneously are called Redox reactions. In the name Redox, the term 'red' stands for reduction and 'ox' stands for oxidation.

Q.1. Elaborate the Equation:

Ans. SO₂ is reduced to sulphur, so it is an oxidising agent. H₂S is oxidized to sulphur, so it is a reducing agent.

It should be noted that substance which undergoes oxidation acts as reducing agent whereas the substance undergoing reduction acts as oxidising agent.
There is another concept of oxidation and reduction in terms of metals and nonmetals:
(i) The addition of nonmetallic elements (or removal of metallic elements) is called Oxidation.
(ii) The addition of metallic elements (or removal of nonmetallic elements) is called Reduction.

Q.2. A shiny brown colored element 'X' on heating in the air becomes black in colour. Name the element 'X' and the black coloured compound formed.
Ans. An element of heating in air changes in its oxide. The brownish element which forms black oxide is copper. 
Name of the element is Copper (Cu)
Name of black compound: Copper(II) oxide, (CuO)
Reaction: 


6. ENDOTHERMIC AND EXOTHERMIC REACTIONS
(i) Endothermic Reactions: The endothermic process is a term that describes a reaction where the system absorbs the energy from its surrounding in the form of heat.
Examples: Photosynthesis, Evaporating liquids, Melting ice, Dry ice, Alkanes cracking, Thermal decomposition, Ammonium chloride in water and much more.

(ii) Exothermic Reactions: The exothermic reaction is the opposite of an endothermic reaction. It releases energy by light or heat to its surrounding.
Examples: Neutralization, burning a substance, reactions of fuels, deposition of dry ice, respiration, solution of sulfuric acid into water and much more.

Difference Between Endothermic and Exothermic Reactions:

ELECTRONIC CONCEPT FOR OXIDATION AND REDUCTION
Oxidation: The loss of an electron by atoms or ions is called Oxidation.
Atom → Cation + electrons
A → Aⁿ⁺ + ne^-
Atom 'A' loses n electrons to become a positively charged ion Aⁿ⁺. It is called a Cation.

Reduction: The gain of an electron by an atom or ion is called reduction.
B + ne^- → B^n-
The atom B gains n' electrons to become negatively charged ion B^n-, it is called Anion.

Redox reactions: Oxidation and reduction reactions occur simultaneously and are called Redox reactions. Only oxidation or only reduction is called half-reaction.
A + B → A⁺B^- → AB  
Example: Na + Cl → NaCl
In this process, sodium loses one electron and oxidised to Na⁺, chlorine gains this electron and is reduced to Cl^-.
Na → Na⁺ + e^- (Loss of an electron is oxidation)
Cl + e^- → Cl^- (Gain of an electron is reduction)
The above two reactions are called half-reactions. These half reactions when added results in complete reaction.

EFFECTS OF OXIDATION REACTIONS IN EVERYDAY LIFE
Oxidation has a damaging effect on metals as well as on food. The damaging effect of oxidation on metals is studied as corrosion and that on food is studied as rancidity.

Thus there are two common effects of oxidation reactions areas:
(i) Corrosion of metals
(ii) Rancidity of food

(i) Corrosion of metals: Corrosion is the process of deterioration of metals as a result of its reaction with air, moisture, and acids (Present in the environment) surrounding it.

Fig: Different conditions of an iron nail rusting

The corrosion causes damage to buildings, bridges, ships and many other articles specially made of iron.
Rust: Iron corrodes readily when exposed to moisture and gets covered with a brown flaky substance called rust.  

Fig: Rusted Chains

This is called rusting of iron, Rust is a hydrated iron (III) oxide Fe₂O₃· 2H₂O.

Rusting of iron takes place under the following conditions:
(a) Presence of air (or oxygen)
(b) Presence of water (or moisture)

It has been observed that:
(i) Presence of impurities in the metal speeds up the rusting process. Pure iron does not rust.
(ii) Presence of electrolytes in water also speeds up the process of rusting
(iii) The position of the metal in the Electrochemical series determines the extent of corrosion. More the reactivity of the metal, there will be more possibility of the metal getting corroded.

Prevention from Rusting:
Methods used to prevent Rusting of Iron are as follows:

• Alloying: Since Rusting of Iron is a chemical process that happens because the metal is attaining a more stable chemical state, alloying (mixing) the iron with other stable metals or alloys can slow down the process of rusting.

Fig: Alloyed Bar

• Galvanizing: Galvanizing a metal object means to coat the surface of that object with a layer of metallic zinc. Also, it is an inexpensive procedure. In conclusion, it will provide it with protection against rusting.

Fig: Galvanizing

• Coating and Painting: Coating the surface of a metal object with a layer of either Paint or Varnish will break the contact between the surface and atmospheric oxygen making it consequently immune to rusting.

Fig: Coating and Painting

• Humidity Control: Controlling the humidity of the environment is also a solution. Therefore, the chances of the metal object rusting will reduce.

Other examples of Corrosion are:

• Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and acquires a green coating of basic copper carbonate.

Fig: Copper corrosion

• Silver articles become black after sometimes when exposed to air because it reacts with sulfur to form a coating of silver sulfide.

Fig: Black silver sulfide deposition

• Lead or stainless steel lose their luster due to corrosion. Unreactive metals such as gold, platinum, palladium, titanium, etc. do not corrode. 

(ii) Rancidity of Food: Fresh foods containing fats and oils smell and taste pleasant but when they remains exposed in air for a long time it's smell and taste changes to unpleasant. It is said that food has become rancid.
OR
It is due to the oxidation of fats and oils, butter, ghee, boiled rice, etc, after prolonged exposure to air i.e. The condition produced by the aerial oxidation of fats and oils in foods marked by unpleasant smell and taste is called rancidity.

Fig: Rancidity in food products

Prevention of rancidity:
(i) Rancidity can be prevented by adding antioxidants to foods containing fats and oils. Antioxidants are reducing agents so when they are added to food it does not get oxidised easily and hence do not turn rancid.
The two common antioxidants are:

BHA (Butylated Hydroxy Anisole)
BHT (Butylated Hydroxy Toluene)

Vitamin-E and vitamin-C (ascorbic acid) are the two antioxidants occurring in natural fats.
(ii) Rancidity can be prevented by packaging fat and oil-containing foods in nitrogen gas.
(iii) It can be retarded by keeping food in the refrigerator.
(iv) It can also be retarded by storing food in airtight containers.
(v) It can be retarded by storing foods away from light.